1. Periodic Relationships Among the Elements Chapter 8 Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

2. How do we arrange and categorize atoms? 19 th Century – Did not know about protons, neutrons and electrons  Try organizing elements by their mass 1864 John Newlands noticed when atoms are arranged by mass we see similar properties every 8 elements  Law of Octaves

3. 1869 Dmitri Mendeleev & Lother Meyer independently proposed more extensive tabulation based on regular, periodic recurrence of properties  Grouped elements more accurately  Predicted that other elements should exist Mendeleev’s periodic table had 66 known elements By 1900 30 more elements were added

4. 8.1 When the Elements Were Discovered

5. What happens if we arrange elements solely by atomic mass? Ar = 39.95 amu > K = 39.10 amu Co = 58.93 amu > Ni = 58.69 amu Te = 127.6 amu > I = 126.9 amu Predict K is a noble gas and Ar is an alkali metal !! 1913 Henry Moseley proposed the idea of an atomic number and predicted the frequency (ν) of x-rays emitted from elements based on atomic number (Z)  ν = a (Z-b) where a & b are constants (same for all elements)

6. Atomic number = number of protons = number of electrons (when neutral) => Electron configuration predicts chemical properties => Look at outermost ground-state electron configurations => valence electrons Li 1s 2 2s 1 Na 1s 2 2s 1 2p 6 3s 1 s1s1 Electron involved in chemical bonding

7. 8.2 ns 1 ns 2 ns 2 np 1 ns 2 np 2 ns 2 np 3 ns 2 np 4 ns 2 np 5 ns 2 np 6 d1d1 d5d5 d 10 4f 5f Ground State Electron Configurations of the Elements

8. Categorize elements by the type of subshell being filled  Representative elements  main group elements  Groups 1A  7A  incompletely filled s & p orbitals  Noble gases  Group 8A  filled s & p orbitals  Transition metals  Group 1B, 3B  8B  incompletely filled d orbitals  Group 2B - Zn, Cd, Hg  filled d orbitals  do not fit transition metals or representative elements  no special name  Lanthanides & Actinides  incompletely filled f orbitals

9. 8.2 Classification of the Elements

10. Electron Configurations of Ions Cations - Positive charge => Remove electrons Li 1s 2 2s 1 3 electrons Li + 1s 2 2 electrons Anions - Negative charge => Add electrons O 1s 2 2s 1 2p 4 8 electrons O 2- 1s 2 2s 1 2p 6 10 electrons

11. Electron Configurations of Cations and Anions Na [Ne]3s 1 Na + [Ne] Ca [Ar]4s 2 Ca 2+ [Ar] Al [Ne]3s 2 3p 1 Al 3+ [Ne] Atoms lose electrons so that cation has a noble-gas outer electron configuration. H 1s 1 H - 1s 2 or [He] F 1s 2 2s 2 2p 5 F - 1s 2 2s 2 2p 6 or [Ne] O 1s 2 2s 2 2p 4 O 2- 1s 2 2s 2 2p 6 or [Ne] N 1s 2 2s 2 2p 3 N 3- 1s 2 2s 2 2p 6 or [Ne] Atoms gain electrons so that anion has a noble-gas outer electron configuration. Of Representative Elements 8.2

12. +1+2+3 -2-3 Cations and Anions Of Representative Elements 8.2

13. Na + : [Ne]Al 3+ : [Ne]F - : 1s 2 2s 2 2p 6 or [Ne] O 2- : 1s 2 2s 2 2p 6 or [Ne]N 3- : 1s 2 2s 2 2p 6 or [Ne] Na +, Al 3+, F -, O 2-, and N 3- are all isoelectronic with Ne What neutral atom is isoelectronic with H - ? H - : 1s 2 same electron configuration as He 8.2

14. Electron Configurations of Cations of Transition Metals 8.2 When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n – 1)d orbitals. Fe: [Ar]4s 2 3d 6 Fe 2+ : [Ar]4s 0 3d 6 or [Ar]3d 6 Fe 3+ : [Ar]4s 0 3d 5 or [Ar]3d 5 Mn: [Ar]4s 2 3d 5 Mn 2+ : [Ar]4s 0 3d 5 or [Ar]3d 5

15. Effective nuclear charge (Z eff ) is the “positive charge” felt by an electron. Na Mg Al Si 11 12 13 14 10 1 2 3 4 186 160 143 132 Z eff Core Z Radius Z eff = Z -  0 <  < Z (  = shielding constant) Z eff  Z – number of inner or core electrons Inner orbitals shield outer orbitals. How many protons will the electron feel? 1s orbital will shield 2s orbital, but 2s orbital will not shield 1s orbital. For a Period as Z eff increases radius decreases.

16. Effective Nuclear Charge (Z eff ) 8.3 increasing Z eff

17. 8.3 2r Atomic radius = ½ distance between two nuclei of adjacent atoms Metals Non-Metals that occur as diatomic molecules

18. Increasing nuclear charge Increasing effective charge Increasing shell number

19. Atomic Radii 8.3

20. Cation is always smaller than atom from which it is formed. Anion is always larger than atom from which it is formed. What about the size of ions ? Generally empty all valence electrons. Add electrons to fill orbital. Get more repulsive forces between electrons.

21. 8.3 Shift smaller Shift larger

22. 8.3 Metals Non- Metals

23. Energy of orbitals in a multi-electron atom Energy depends on n and l E = 0

24. Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state. I 1 + X (g) X + (g) + e - I 2 + X (g) X 2 + (g) + e - I 3 + X (g) X 3 + (g) + e - I 1 first ionization energy I 2 second ionization energy I 3 third ionization energy I 1 < I 2 < I 3 It gets harder and harder to remove electrons !!

25. 8.4 Where do we see the highest ionization energies?

26. Filled n=1 shell Filled n=2 shell Filled n=3 shell Filled n=4 shell Filled n=5 shell 8.4 Trend to smaller I 1 due to increasing distance from nucleus

27. General Trend in First Ionization Energies 8.4 Increasing First Ionization Energy

28. Electron affinity is the negative of the energy change that occurs when an electron is accepted by an atom in the gaseous state to form an anion. X (g) + e - X - (g) 8.5 F (g) + e - X - (g) O (g) + e - O - (g)  H = -328 kJ/mol EA = +328 kJ/mol  H = -141 kJ/mol EA = +141 kJ/mol Reverse the sign

29. 8.5

30. Orbitals that are almost filled want an electron Noble gases do not want to accept an electron

31. 8.6 Charge density is similar along the diagonal => expect similar properties Happens most with lighter elements.

32. Hydrogen (1s 1 ) Where should we put it in the periodic table? - Hydrogen is in a class by itself - has 1 electron is s orbital like alkali metals => can form H+ can also gain an electron to form H- => hydride ion => resembles halogens

33. Ionic hydrides react with water to produce H 2 gas + metal hydroxides NaH (s) + H 2 O (l) NaOH (aq) + H 2 (g) Hydrogen Displacement CaH 2 (s) + 2 H 2 O (l) Ca(OH) 2 (aq) + 2 H 2 (g) 2 H 2 (g) + O 2 (g) 2 H 2 O (l) Hydrogen can also burn in air

34. Group 1A Elements (ns 1, n  2) M M +1 + 1e - 2M (s) + 2H 2 O (l) 2MOH (aq) + H 2(g) 4M (s) + O 2(g) 2M 2 O (s) Increasing reactivity 8.6 Low first ionization energies Metals lose their shiny appearance to form oxides.

35. 4 Li (s) + O 2 (g) 2 Li 2 O (s) Some alkali metals can form peroxides (O 2 2- ion) 2 Na (s) + O 2 (g) Na 2 O 2 (s) K, Rb & Cs can form superoxides (O 2 - ion) K (s) + O 2 (g) KO 2 (s) As we move down the periodic table we can for other oxide forms.

36. Group 2A Elements (ns 2, n  2) M M +2 + 2e - Be (s) + 2H 2 O (l) No Reaction Increasing reactivity Mg (s) + 2H 2 O (g) Mg(OH) 2(aq) + H 2(g) M (s) + 2H 2 O (l) M(OH) 2(aq) + H 2(g) M = Ca, Sr, or Ba Less reactive than alkali metals. Form divalent ions Reacts slowly with steam Much more reactive cold Increasing metallic character

37. Increasing metallic character as you go down the group. BeH 2, BeCl 2, and MgH 2 are more molecular than ionic in nature Strontium - 90 is radioactive - Major product in atomic bomb explosions - Since Sr 2+ and Ca 2+ are chemically similar, Sr 2+ can replace Ca 2+ in our bones => high energy radiation => can cause anemia, leukemia and other chronic illnesses

38. Group 3A Elements (ns 2 np 1, n  2) 8.6 4Al (s) + 3O 2(g) 2Al 2 O 3(s) 2Al (s) + 6H + (aq) 2Al 3+ (aq) + 3H 2(g) Boron is a metalloid and the rest are metals Boron does not react with O 2 (g) & water Forms unreactive oxide coating Al only forms 3+ ions Other 3A elements form + & 3+ ions Increasing metallic character

39. Group 4A Elements (ns 2 np 2, n  2) 8.6 Sn (s) + 2H + (aq) Sn 2+ (aq) + H 2 (g) Pb (s) + 2H + (aq) Pb 2+ (aq) + H 2 (g) Increasing metallic character non-metal metalloids metals C, Si & Ge do not form ionic compounds Sn & Pb will not react with water, but will react with HCl

40. Group 4A elements form compounds with 2+ & 4+ oxidation states 4+ is generally the most stable  CO 2 is more stable than CO  SiO 2 is very stable and SiO does not exist under normal conditions For Pb the stability trend is reversed 2+ is more stable than 4+ Pb [Xe] 6s 2 4f 14 5d 10 6p 2 Lose only 6p 2 electrons => Pb 2+ [Xe] 6s 2 4f 14 5d 10

41. Group 5A Elements (ns 2 np 3, n  2) non-metals metalloids metal Increasing metallic character Larger variation in properties Nitrogen  forms diatomic gas - N 2  forms oxides - NO, N 2 O, NO 2, N 2 O 4 & N 2 O 5 => NO x  all oxides are gases except N 2 O 5 (s)  accepts 3 electrons to become isoelectric with Ne  compounds w/ metals are ionic – Li 3 N & Mg 3 N 2

42. N 2 O 5(s) + H 2 O (l) 2 HNO 3(aq) P 4 O 10(s) + 6 H 2 O (l) 4 H 3 PO 4(aq) Phosphorus  exists as P 4 molecules  forms solid oxides as P 4 O 6 & P 4 O 10 N & P form important oxoacids Arsenic Antimony Bismuth Form extensive 3-d networks

43. Group 6A Elements (ns 2 np 4, n  2) 8.6 non-metals metalloids No metals in this group Oxygen form diatomic gas – O 2 Sulfur forms molecular S 8 Selenium forms molecular Se 8 Tellurium & Polonium – form extensive 3-d networks Polonium is radioactive

44. SO 3(g) + H 2 O (l) H 2 SO 4(aq) Oxygen tends to form O 2- ion in many compounds S 2-, Se 2-, and Te 2- ions are common S forms important oxides – SO 2, SO 3 => SO x

45. Group 7A Elements (ns 2 np 5, n  2) X + 1e - X - 1 X 2(g) + H 2(g) 2HX (g) Increasing reactivity 8.6 All halogens are non-metals => X 2 => Want to gain 1 electron => halide ions => acids Halogens are very reactive 2 F 2 (g) + H 2 O (l) 4 HF (aq) + O 2 (g)

46. Group 8A Elements (ns 2 np 6, n  2) 8.6 Completely filled ns and np subshells. Highest ionization energy of all elements. No tendency to accept extra electrons. Occur as monatomic gases. Also called inert gases.

47. 1963 – Neil Bartlett made the first Xe compounds Xe (g) + PtF 6 (g) XePtF 6 (s) Since then XeF 4, XeO 3, XeO 4, XeOF 4, KrF 2 have been prepared.

48. Properties of Oxides Across a Period basicacidic 8.6